chemical changes & structure higher chemistry

KHS 2014 - Cheviot Learning Community page 1

Chemical Changes & Structure Topic 2 - Reaction Rates

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Higher Chemistry

Topic 2:

Reaction RatesStudent:

1. Concentration and Collisions

2. Particle Size and Collisions3. Temperature and Collisions

Check Test Score: / 8 Home Practice Score: / 10

3.1CollisionTheory

1. Catalytic Converters

2. Heterogeneous Catalysts3. Homogeneous Catalysts


3.2CatalystsAt Work

1. Considering the Investigation

2. The Experimental Work3. Processing the Results


3.3ConcentrationAnd Rate

1. Considering the Investigation

2. The Experimental Work3. Processing the Results


3.4TemperatureAnd Rate

1. Progreass of a Reaction

2. Calculating the Rate3. Comparing Reaction Progress


3.5ReactionProgress

1. Energy During Collision

2. Potential Energy Diagrams3. Catalytic Effect


3.6ActivationEnergy

1. Distribution Graphs

2. Concentration and Temperature3. Catalytic Effect


3.7EnergyDistribution

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Consolidation B Score: / 10 Consolidation C Score: / 10

Consolidation D Score: / 10

ConsolidationWork

End-of-SectionAssessment

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Lesson Activities Done Checked



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2.1 Collision Theory

Concentration & Collisions

This first lesson looks at how the effect of concentration, particle size and temperature on reaction rates can be explained in terms of collisions between reactant particles.

This activity looks at how the collision theory can explain why the concentration of reactants affects the rate of a chemical reaction.

The Collision Theory is simply common sense.

For a reaction to occur the particles of two reactants have to make contact.; they must collide.

The reactant particles must collide with enough force to break the existing bonds in their particles.

In the reaction between marble and hydrochloric acid, the hydrogen ions collide with the surface of the marble.

In more concentrated acid, there are more hydrogen ions per cm3, so there will be more collisions per second.

Not every collision will have enough energy, but there will be more successfulcollisions per second than before,so the rate of the reaction will be increased.

1M acid

4M acid

The Collision Theory states that substances can only react with each other if ( i) their particles collide with each other, and (ii) the collisions have enough energy.

Gentle collisions just result in reactant particles bouncing apart unchanged.

Breaking bonds requires energy, and that energy comes from the force of the collision.



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large marble particleIn the reaction between marble and hydrochloric acid, only collisions on the surface of the marble can lead to reactions

In the large marble particle, only 28 marbleions (CO3

2-) are on the surface.

The mass of the small marble pieces is the same as the large marble, and so, therefore,are the number of marble ions,

This time, there are 48 marble ions on thesurface of the small marble lumps.

With a larger surface area available, there will be more collisions per second which will result in a faster reaction.

small marble particles

Temperature & Collisions This activity looks at how the collision theory can explain why the temperature of the reactants affects the rate of a chemical recation.

Increasing the temperature causes the reactant particles to move faster. This means that they collide more often and also that they collide with greater kinetic energy. In the reaction between marble and hydrochloric acid, it willbe the hydrogen ions that are most affected by an increase intemperature.ons

At the higher temperature, there will bemore collision per second and this willresult in a faster reaction.

More importantly, at the higher temperature,the collisions will be more ‘violent’ ,(higherenergy), making a reaction more likely thanwith collisions at the lower temperature.

With more collisions per second taking place, and with a larger proportion of these collisions leading to a reaction, there will often be a dramatic increase in the reaction rate.

With many reactions, an increase of only 10°C is often enough to double the rate of the reaction.

high temperature

low temperature

Particle Size & Collisions This activity looks at how the collision theory can explain why the particle size of a solid reactant affects the rate of a chemical reaction.

The word particle, in this context, refers to the form of the solid; lump, grains, powder etc



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2.2 Catalysts at Work

Catalytic Converters

This lesson considers catalysts at work in different kinds of chemical reactiions.

This activity looks at how the catalytic converter in a car exhaust system converts polluting gases into non-polluting substances.

Converters are needed to change harmful gases into less harmful gases. Most have two sections:-

Section 1 - platinum & rhodium catalysts

+ → +nitrogen monoxide nitrogen dioxide nitrogen oxygen

Section 1 - platinum & palladium catalysts

+ → carbon dioxide oxygen carbon dioxide

exhaustpipe

silencer

stainless steelhousing

ceramic honeycomb structure covered with (expensive) platinum

and rhodium

unburnt petrol,oilcarbon monoxideoxides of nitrogen

nitrogen, oxygenwater vapour

carbon dioxide

OO C O

O C O C O O

The gas molecules collide with the catalyst and stick to its surface. This is called adsorption.

Weak bonds form between the reactant molecules and the catalyst.

This weakens the original bonds within the molecules.

This makes it much ‘easier’ for the reactant molecules to react to form products.

‘Easier’ means the collisions can be less violent, so a reaction can occur at a lower temperaturethan normal, or will proceed at a faster rate thannormal for that temperature.

OO C O

C OO

CO

O

Once the product has formed, the weak bonds with the catalyst break and the product molecules depart.

The surface of the catalyst is then available to adsorb more reactant molecules and repeat the process.

O C O C O OO O C O



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Heterogeneous Catalysis This activity is about examples of heterogeneous catalysis.

OO C O

O C O C O OPb Pb Pb Pb

Cars with catalytic converters must be able to run on unleaded petrol.

The lead atoms can form strong bondson the surface of the catalyst and so prevent the adsorption of reactant molecules.

This effect is known as poisoning the catalyst. Catalysts of this type are often referred to as surface catalysts and benefit from a large surface area, so are often used as powders or sprayed onto a honeycomb structure. (This can also make a small amount of, very expensive, transition metal go much further).

In heterogeneous catalysis the catalyst is in a different state from the reactants.

EXAMPLES OF HETEROGENEOUS CATALYSIS

Process Reactants Catalyst Important Product

During the industrial cracking of large oil molecules the powdered aluminosilicate catalyst becomes coated with carbon (soot). This is another example of catalyst poisoning. It can-not be avoided, but the powder is continually removed and the soot burned off to allow the catalyst to be used again. This is known as Regeneration.



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Homogeneous Catalysis This activity is about examples of homogeneous catalysis.

In homogeneous catalysis the catalyst is in the same state as the reactants, (usually that of solution).

In ordinary life, many reactions take place in the solutions inside your body cells and are catalysed by natural catalysts called enzymes.

Enzymes are biological catalysts. They are complex protein molecules which speed up specific chemical changes by homogeneous catalysts.

ENZYMES AT WORK

Industry Enzyme or Change Achieved its Source

invertase sugar → +

winemaking →

yoghurtmaking milk →

yeast → beer

Hydrogen peroxide solution and Rochelle salt solution react together to produce bubbles of oxygen gas. The reaction is extremely slow at room temperature, but at 50°C the bubbles of oxygen are noticeable.

Pink cobalt chloride solution (Co2+(aq)), when added to the mixture, begins to turn green

(Co3+(aq)). At the same time, the speed of the reaction increases rapidly.

As the reaction subsides, the pink colour of Co2+(aq) returns. Like all true catalysts, the cobalt

chloride is unchanged at the end of the reaction,



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2.3 Concentration and RateThe purpose of this lesson is to investigate in more detail the way that the concentration of a reactant affects the rate of a chemical reaction.

PPA* State the aim of the experiment

Procedure* Describe how the concentration of the potassium iodide solution was varied

* State two factors which had to be constant in the experiments.

* How was the rate of the reaction determined ?

Results

* Present Your results in a table:



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* Draw a graph to illustrate your results

Conclusion* State the conclusion of the experiment

Evaluation* The procedure adopted to determine the reaction time was particularly effective and accurate. Suggest why this is the case.

* Apart from allowing us to take the volume of potassium iodide as a measure of its concentration, what was the purpose of keeping the total volume of the reaction mixture the same in each experiment?



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2.4 Temperature and RateThe purpose of this lesson is to investigate in more detail the way that the temperature of the reactants affects the rate of a chemical reaction.

PPA* State the aim of the experiment

Procedure* State two factors which had to be kept constant in the experiments.

* How was the rate of the reaction determined ?

Results* Present your results in a tabular form:



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* Draw a graph to illustrate your results

Conclusion* State the conclusion of the experiment

Evaluation* Why did the beaker in each experiment have to be dry?

* Suggest why it might be difficult to get an accurate reaction time if the experiment was carried out at room temperature



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2.5 Reaction ProgressThis lesson topic deals with some ways of following the progress of a chemical reaction.

Progress of a Reaction The aim of the following experiment is to follow the progress of a reaction by recording the volume of gas produced at regular intervals.

calcium + hydrochloric → calcium + carbon + water carbonate acid chloride dioxide

CaCO3 + 2 HCl → CaCl2 + CO2 + H2O

Time (min)

0

2

4

6

8

10



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A shallow slope - many reactions are slow to get started

B steep slope - fast reaction rate

C shallow slope - reaction starts to slow down as chemicals are used up (their concentrations fall)

D level slope - reaction has stopped. One of the chemicals has been used up completely

Quantity being measured

Time

A

B

C

D



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Calculating the Rate This activity examines how the rate of a reaction can be calculated from a progress graph.

Rate of reaction is the change in quantity of a reactant or product per unit of time.

average rate =change in time

change in quantity

20 40 60 80 100 120 140Time ( s )

0

10

20

30

40

50

60

70

80

90

100

110

120

Volu

me

of h

ydro

gen

( cm

3 )

The unit used for rate depends on the quantity of the reactant/product that is being measured, and the time scale for the reaction. e.g weight loss (electrical balance) grammes g/s , g/min, g/hour gas volume (syringe) ml or cm3 cm3/s etc. concentration (colourimeter) moles/litre moles/l/s etc.

The reaction between sulphuric acid and magnesium produces hydrogen gas. The progress of the reaction can be monitored by measuring the volume of gas produced. The Progress Graph, below, can be used to calculate the rate of this reaction at different stages.



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Time interval Change in volume Average rate ( s ) ( cm3 ) ( cm3 s-1 )

0 — 20

20 — 40

40 — 60

60 — 80

80 — 100

100 — 120

120 — 140

20 40 60 80 100 120 140Time ( s )

The rate will be at a maximum near the beginning of the reaction, (when the concentrations of the reactants are at their highest level), will usually drop quite steadily (as the reactant concentrations decrease) and will eventually reach zero (once one of the reactants is used up completely.)



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Comparing Reaction Progress The purpose of this activity is to compare progress graphs for the same reaction taking place under different conditions

Example 1.

5 10 15Time ( min )

0

40

80

120

20

60

100

Volume of gas ( cm3 )lower

temperature

lowertemperature

5 10 15Time ( min )

194

198

200

193

196

199

Mass of flask & contents ( g )

lowerconcentration

lowerconcentration

Example 3.

5 10 150

40

80

120

20

60

100

Volume of gas ( cm3 )

uncatalysedreaction

catalysedreaction

Example 2.

The faster reaction will produce more gas over the same time interval:- the slope will be steeper.

The faster reaction will finish first.

Both reactions have used the same mass of zinc, with the sameparticle size, with the same volume and concentration of hydrochloric acid. The final volumes of hydrogen gas will, therefore, be the same . The faster reaction was at a higher/lower temperature

The faster reaction will lose more mass over the same time interval:- the slope will be steeper.

The faster reaction will finish first.

Both reactions have used the same mass of marble, with the same particle size, with the same volume of hydrochloric acid at the same temperature. The final masses of the flask & contents will, therefore, be the same . The faster reaction was at a higher/lower concentration

The catalysed reaction will bethe faster reaction and will produce more gas over the same time interval:- the slope will be steeper.

The catalysed reaction will finish first.

Both reactions have used the same mass of zinc, with the same particle size, with the same volume and concentration of sulphuric acid at the same temperature, so the final volume of gas will be the same.



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Reaction progress

Potential Energy

2.6 Activation EnergyThis lesson explores the idea of activation energy and how it affects chemical reactions.

Energy During Collision This activity looks at how the chemical energy of a reaction mixture changes as the reactant molecules collide and change to form the product molecules.

Imagine a very simple reaction between two diatomic molecules:

A — A + B — B → 2 A — B

→ →

At the start of the reaction the reactant molecules have kinetic energy and potential energy.

The potential energy is mainly energy stored within their bonds and can also be called chemical energy.

As the reactant molecules begin to collide, strong forces of repulsion build up as their outer electrons repel each other.

The molecules begin to slow down, and their kinetic energy is converted into potential energy.

As the molecules crunch into each other and stop, their kinetic energy is now entirely changed to potential energy.

This extra energy can be enough to break existing bonds and allow new bonds to form.

If new bonds have formed, then as the product molecules spring apart potential energy will change back to kinetic energy.

The bonds in the product molecules will be different from the original reactant molecules.

The products will have a different amount of potential energy from the reactants :-energy will have been lost (exothermic reaction) or energy will have been gained (endothermic reaction).



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Reaction progress

Ea

activated complex

Potential energy ( kJ mol-1 )

0

10

20

30

40

50

Potential Energy Diagrams This activity takes a further look at the potential energy diagrams for chemical reactions

During collisions, kinetic energy is changed to potential energy and stored within the bonds of the reactants

If this energy is large enough, then the reactant molecules will break apart and go on to form products. In between, an unstable arrangement of atoms called the activated complex will exist.

If there is not enough energy, the reactant molecules will separate and proceed unchanged:- no reaction

In a sense there is an energy barrier which has to be overcome before reactant molecules can successfully change into products. The extra potential energy that reactants need before they can react is called the Activation Energy, Ea.

Ea = kJ mole-1

An activated complex is the unstable arrangement of atoms formed at the maximum of the potential energy barrier.

Activation energy is the additional potential energy which has to be achieved by colliding molecules to form an activated complex.

Ea = P.E(activated complex) — P.E(reactants)

Reaction progress


0

20

40

60

80

100

Reaction progress


0

100

200

300

400

500

Ea = 70 kJ mole-1 Ea = 200 kJ mole-1



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Potential Energy Diagrams This activity explains the effect of a catalyst in terms of activation energy.

Reaction progress

Potential Energy

A catalyst provides a different reaction pathway:- a different series of collisions which will still change the reactants into products.

A catalyst lowers the activation energy :- the alternative route is ‘easier’, meaning that less potential energy is needed before bonds break.

Catalysed reactions can be done at lower temperatures :- the reactants need less kinetic energy to ensure a successful collision.

Catalysts speed up a reaction by providing an alternative reaction pathway with a lower activation energy



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2.7 Energy DistributionThis lesson looks at the effect of a distribution of kinetic energies on the rate of a reaction

Distribution Graphs This activity looks at distribution graphs for kinetic energies

Kinetic energy

Number of moleculeswith a given

kinetic energy

0

Molecules are always moving - they have kinetic energy.

At any temperature, however, molecules can have a range of kinetic energies:-

low energy - very few medium energy - most high energy - very few

Similarly, collisions between molecules will have a range of energies

Kinetic energy

Number of collisionswith a given

kinetic energy

0

Somewhere along the x-axis there will be the value for the minimum kinetic energy needed for a collision to lead to a successful reaction - the activation energy.

All collisions with this energy, or above, will lead to a reaction. Shading this area on the graph gives an indication of the number of successful collisions.

The larger this number, the faster the reaction

The normal method for increasing the kinetic energy of reacting particles is to increase the temperature - add heat energy. You should be aware, however, that many reactants are capable of taking in energy in other forms:-

e.g. light energy -photosynthesis, sunlight provides the energy needed to speed up the reaction between CO2 and H2O.

- photography, light provides the energy to help convert silver ions back into silver atoms to produce the dark patches on a film negative

- substitution, the reaction between alkanes and bromine requires sunlight

- H2 / Cl2 the slow reaction between H2 and Cl2 can be speeded up (made explosive) using a UV camera flash



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Concentration & Temperature

This activity looks at how increasing the concentration or increasing the temperature affects an energy distribution graph.

Kinetic energy


kinetic energy

0Kinetic energy


kinetic energy

0Kinetic energy


kinetic energy

0

low concentration increased concentration

Increasing the concentration simply increases the number of collisions across all possible values of kinetic energy. For example, doubling the concentration doubles the height of the graph everywhere.

This, of course, increases the number of collisions with activation energy (or above) which increases the number of successful collisions and increases the rate of the reaction.

Kinetic energy


kinetic energy

0Kinetic energy


kinetic energy

0

low temperature increased temperature

When you use a thermometer to measure the temperature of a substance you are really measuring the average kinetic energy of its particles. A few particles have below average energy, most have close to average while a few have above average energy.

Increasing the temperature causes all the molecules to move faster, to have increased kinetic energy.

The effect on the graph is to spread it out to higher values of kinetic energy. There will always be a few particles with well below average energy so the graph still starts at zero.

There will be more particles at the high end of the energy scale whilst there will have to be a similar decrease in particles in the middle. The height of the graph will be lower.



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A small shift to the right can bring a very large number of particles over the activation energy ‘barrier’. A small increase in temperature can often have a dramatic effect on the rate of a reaction. Often a 10°C rise in temperature is enough to double the rate of reaction.

Catalytic Effect This activity explains how a catalyst speeds up a reaction without having any effect on the energy distribution of the molecules

uncatalysed catalysed

Adding a catalyst has no effect on the number of collisions or on the kinetic energy of these collisions.

A catalyst provides an alternative route with a lower activation energy. Many collisions that, previously, had been lacking the energy needed to lead to a reaction will now prove to be successful.


Kinetic energy


kinetic energy

0Kinetic energy


kinetic energy

0



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UNIT 1. Chemical Change & Structure

Plot graphs of changes against time; slope = rate

Average rate = (change in quantity) / time

Rate = 1 / time

concentration↑ number of collisions↑particle size↓ surface area↑ number of collisions↑

e.g photosynthesis photography

heterogeneous - different state from reactantshomogeneous - same state as reactants

Topic 2: Reaction Rates

Following the course of a reaction1. Reactions can be followed by measuring changes in concentration, mass and volume of reactants or products.

2. The average rate of a reaction, or part of a reaction, can be calcukated from initial and final quantities and the time interval

3. The rate of a reaction, or part of a reaction, is proportional to the reciprical of the time taken.

Factors affecting rate4. The rates of reactions are affected by changes in concentration, particle size and temperature.

5. The collision theory can be used to explain the effects of concentration and particle size (surface area) on reaction rates

6. Temperature is a measure of the average kinetic energy of the particles of a substance.

7. The activation energy is the minimum kinetic energy required by colliding particles before reaction will occur.

8. Energy distribution diagrams can be used to explain the effect of changing temperature on the kinetic energies of particles.

9. The effect of temperature on reaction rate can be explained in terms of an increase in the number of particles with energy greater than the activation energy.

10. With some chemical reactions light can be used to increase the number of particles with energy greater than the activation energy.

Catalysts11. Catalysts can be classified as either heterogeneous or homogeneous.

12. Catalysts are used in many industrial processes.

number of molecules with a

given kinetic energy

kinetic energy

T1



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13. Heterogeneous catalysts work by the adsorption of reactant molecules.

14. The surface activity of a catalyst can be reduced by poisoning.

15. Impurities in reactants mean that industrial catalysts have to be regenerated or renewed.

16. Catalytic convertors are fitted to cars to catalyse the conversion of poisonous carbon monoxide and oxides of nitrogen to carbon dioxide and nitrogen.

17. Cars with catalytic convertors only use “lead- free” petrol to prevent poisoning of the catalyst.

18. Enzymes catalyse the chemical reactions which take place in the living cells of plants and animals.

19. Enzymes are used in many industrial processes.

Potential energy diagrams20. A potential energy diagram can be used to show the energy pathway for a reaction.

21. The activated complex is an unstable arrangement of atoms formed at the maximum of the potential energy barrier, during a reaction.

22. The activation energy is the energy required by colliding molecules to form an activated complex.

23. The activation energy can be calculated from potential energy diagrams.

24. The effect of a catalyst can be explained in terms of alternative reaction pathways with lower activation energy.

25. A potential energy diagram can be used to show the effect of a catalyst on activation energy.

Often called ‘surface’ catalysts

Impurities adsorbed strongly onto surface.

Made predictions.

e.g 2CO + 2NO → 2CO2 + N2

Progress of reaction

Potential energy

EAExothermic

Endothermic

Ea = maximum energy - energy of reactants

chemical changes & structure higher chemistry

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